Calcium - Element information, properties and uses | Periodic Table Chemistry in its element: calcium (Promo) You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry. (End promo) Chris Smith Hello, welcome to this week's Chemistry in its Element, I'm Chris Smith. This week it's the turn of the element that gives us cement, plaster of Paris, our own bones, hard teeth and hard water. Karen Faulds Milk, cheese, yogurt, spinach, almonds. What element do they all have in common? It's calcium of course! But whilst most off us immediately think of food when someone mentions calcium (and I personally hold the old milk TV adverts accountable for this), it actually has a far bigger role in our lives than that. Calcium is all around us. The average human contains approximately 1kg of calcium, of which 99% is stored in our bones. It is the 5th most abundant element in the earth's crust, occurring widely as calcium carbonate which is more commonly known as limestone. It is also the fifth most abundant dissolved ion in seawater.

Calcium was named after the Latin term calx meaning lime, and is a reactive silvery metallic element found in Group 2 of the periodic table. It was first isolated in 1808 in England when Sir Humphry Davy electrolyzed a mixture of lime and mercuric oxide. Today we obtain calcium through the electrolysis of a fused salt such as calcium chloride. Once exposed to air, elemental calcium rapidly forms a grey-white oxide and nitride coating. Unlike magnesium, calcium is quite difficult to ignite, but once lit, it burns with a brilliant high-intensity red flame. The compounds of calcium are however much more useful than the element itself. Literature dating back to 975 AD shows that plaster of Paris (which is calcium sulphate) was used even then for setting broken bones. Calcium oxide (also known as lime or quicklime) is a major component of mortar and cement. The production of cement using calcium oxide has long been known; it was used by the Romans and also the Egyptians who built the Great Pyramid of Giza and Tutankhamen's tomb.

Calcium fluoride is also well known for being insoluble and transparent over a wide range of wavelengths, making it useful for making cells and windows for infrared and ultraviolet spectrometers. Our drinking water also contains calcium ions - more so in so called hard water areas. Hard water is the term used for water with a high proportion of calcium and magnesium (2 plus) ions. The calcium usually enters the water as it flows past either calcium carbonate, from limestone and chalk, or calcium sulfate, from other mineral deposits. Whilst some people do not like the taste, hard water is generally not harmful to your health. Although it does make your kettle furry! Interestingly, the taste of beer (something dear to my heart) seems related to the calcium concentration of the water used, and it is claimed that good beer should have a calcium concentration that is higher than that of hard tap water. Calcium is what is known as an essential element, meaning that it is an element which is absolutely necessary for life processes. Which is what the old milk TV adverts were trying to tell us after all. Calcium is used to produce the minerals contained in bones, shells and teeth through a process called biomineralisation.

Calcium phosphate (also known as hydroxyapatite) is the mineral component of bones and teeth and is a particularly good example of how organisms fabricate 'living' composite materials. Indeed, the different properties (such as stiffness) of bone are produced by varying the amount of organic component, mostly a fibrous protein called collagen, with which hydroxyapatite is associated. The bone in our body functions not only as a structural support, but also as the central Ca store. Thus, during pregnancy, bones tend to be raided for their Ca in a process called demineralisation. Bone does not last forever; a serious medical problem is osteoporosis which is the decalcification of bone. This loss of bone mass which occurs with increasing age makes bones more susceptible to breaking under stress and it occurs mainly in older people, especially women. Calcium ions also play a crucial role in higher organisms as an intracellular messenger. Fluxes of Ca2+ trigger enzyme action in cells in response to receiving a hormonal or electrical signal from elsewhere in the organism. Calcium is also very important in helping blood to clot. When bleeding from a wound suddenly occurs, platelets gather at the wound and attempt to block the blood flow.

Calcium, vitamin K, and a protein called fibrinogen help the platelets to form a clot. If your blood is lacking calcium or one of these other nutrients, it will take longer than normal for your blood to clot. The ability to detect extremely small amounts of an element can be a very useful adaptation for an animal if that element is important to it. For example, hermit crabs, which inhabit second hand shells and change to newer, bigger shells as they grow, have the ability to recognise shells suitable for occupation not only by feeling for them, but apparently also by measuring the minute amount of calcium carbonate that is dissolved in the water around a shell. They can readily distinguish natural shells containing calcium carbonate from calcium-bearing replicas made from calcium sulphate. The concentration of calcium detected by the hermit crab is in the order of 4ppm or less, which is amazingly low. So from strong teeth and bones, through to good tasting beer and ensuring hermit crabs find their perfect home -you can see that calcium really is an essential element.

Chris Smith Well, I'm very at home with my hard water, and the local beer tastes quite good too, although I do get through quite a few kettles - indeed Russell Hobbs probably owe their buoyant share price just down to me. Well, maybe. That was Strathclyde University's Karen Faulds with the story of Calcium. Next week, if you were an element which one would you be? Pat Bailey If I had to choose a person to represent gold, then I guess it might be an ambitious young stockbroker, a bit flashy, and not great at forming relationships. For helium - an airy-fairy blonde with a bit of a squeaky voice, but with aspirations to join the nobility. And for boron? Well at first glance a boring, middle-aged accountant, maybe wearing brown corduroys and a tweed jacket . but with an unexpected side-to him in his spare time - skydiving, and a member of a highly dubious society that indulges in swapping partners. Chris Smith And you can get the inside story on Boron's swinging antics with Pat Bailey in next week's Chemistry in its Element. I'm Chris Smith, thank you for listening and goodbye.

(Promo) Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com . There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements . (End promo) Calcium, Chemical Element - structure, reaction, water, uses, elements, examples, metal, gas CAL-cee-um Discovery and naming It is impossible to say when humans first knew about or used compounds of calcium. Whenever they used limestone to build a structure, they were using a compound of calcium. Limestone is the common name for calcium carbonate (CaCO 3 ). Whenever humans built a statue or monument out of marble, they were using calcium carbonate in another form. Ancient Egyptians and early Greeks used mortar, a cement-like material that holds stones and bricks together. Early mortar was made by roasting or heating limestone for long periods of time. Water was then mixed with the powder, which would then dry to form a strong bond. Humphry Davy | English chemist H umphry Davy (1788-1829) was a major contributor to the field of electrochemistry. This is the science involving the relation of electricity to chemical changes.

He is best known for discovering calcium, sodium, potassium, magnesium, strontium, and barium. He also discovered nitrous oxide and two types of lamps. Davy grew up in Cornwall, England, in a poor family. His father, who died when Davy was a boy, had lost money in unwise investments, so Davy helped his mother pay off the debts. He disliked being a student, though he liked reading about science. In later life, he said he was happy he did not study too hard because he had more time to think on his own. With no money for further education, the 17-year-old Davy began to work for a surgeon-pharmacist. He also started learning on his own about other subjects that interested him, such as geography, languages, and philosophy. He even wrote poems that later earned him the respect and friendship of William Wordsworth, Samuel Coleridge, and other leading English poets of his time! At 19, Davy read a chemistry book by the famous French scientist Antoine-Laurent Lavoisier (1743-94). That book convinced him to concentrate on chemistry.

For the rest of his life, Davy's career was marked by brilliant scientific explorations in chemistry and electrochemistry. Davy discovered nitrous oxide after testing the effects of hydrogen and carbon dioxide on himself. (He liked to use himself as a human guinea pig!) Nitrous oxide is a gas consisting of nitrogen and oxygen. While studying nitrous oxide gas, he discovered that its effects often made him feel very happy or very sad. The feeling of happiness eventually gave nitrous oxide another name: laughing gas. Most importantly, though, Davy recognized that it could be used as an anesthetic. An anesthetic is a chemical used to dull pain during minor surgery. In 1808, Davy invented the carbon arc lamp. He had proposed using carbon as the electrode material instead of metal. (Electrodes are conductors used to establish electrical contact with a nonmetallic part of a circuit.) With carbon electrodes, he made a strong electric current leap from one electrode to the other. This created an intense white light. Davy's invention marked the beginning of the era of electric light. Arc lamps are still used today.

Using his knowledge of electricity, Davy built a large battery which he used to break down substances most scientists thought were pure elements. In 1807, he discovered the element potassium. He created this by using electrolysis. Electrolysis produces chemical changes by passage of an electric current through an electrolyte. An electrolyte is a nonmetallic electric conductor. Within a week he isolated sodium in a similar way. Then in 1808, he used a slightly modified method to isolate calcium, magnesium, barium, and strontium. Davy was only 29 by the time he had discovered all of these elements! Davy later invented the miner's lamp (now known as the Davy lamp). He learned that methane was the mine gas that caused explosions. But he realized it ignited only at high temperatures. So he designed a lamp in which the flame was surrounded by wire gauze. This reduced the heat and prevented flammable gases from igniting. This made coal-mining safer by reducing the number of explosions in mines. Davy was rewarded by many honors and medals for his discoveries and inventions. He died of a stroke in 1829 at the age of 49.

Another calcium compound used by early civilizations was plaster of paris. Plaster of paris is made by heating gypsum, or calcium sulfate (CaSO 4 ), to remove the water that makes it crystallize. Water was added and it hardened into a brittle, cement-like substance. Until recently, it was most often used to make casts to protect broken bones. However, it has largely been replaced by fiberglass, which is lighter, yet stronger. The first mention of plaster of paris to protect broken bones can be found in a book written by Persian pharmacist Abu Mansur Muwaffaw in about 975 A.D. By the 1700s, chemists had learned a great deal about calcium compounds. They knew that limestone, gypsum, marble, and many other commonly occurring compounds all contain a common element. They called the element calx. That word comes from the Latin term for lime. In 1807, Davy isolated the new element. Davy invented a system for melting compounds of elements that were difficult to separate by usual methods. He passed an electric current through the compound, causing the compounds to break into parts. One of those parts was calx.

He created the name calcium by adding the suffix -ium to calx; -ium is the ending used for almost all metallic elements. Davy was also able to produce free sodium, potassium, strontium, magnesium, and barium. Physical properties Calcium is a fairly soft metal with a shiny silver surface when first cut. The surface quickly becomes dull as calcium reacts with oxygen to form a coating of white or gray calcium oxide. Calcium's melting point is 850°C (1,560°F) and its boiling point is 1,440°C (2,620°F). It has a density of 1.54 grams per cubic centimeter. Chemical properties Calcium is a moderately active element. It reacts readily with oxygen to form calcium oxide (CaO): Calcium reacts with the halogens— fluorine, chlorine, bromine, iodine, and astatine, The halogens are the elements that make up Group 17 (VIIA) of the periodic table. Calcium also reacts readily with cold water, most acids, and most nonmetals, such as sulfur and phosphorus.

For example, calcium reacts with sulfur: Occurrence in nature Calcium is the fifth most common element in the Earth's crust. Its abundance is estimated to be about 3.64 percent. It is also the fifth most abundant element in the human body. Calcium does not occur as a free element in nature. It is much too active and always exists as a compound. The most common calcium compound is calcium carbonate (CaCO 3 ). It occurs as aragonite, calcite, chalk, limestone, marble, and travertine, and in oyster shells and coral. Shellfish build their shells from calcium dissolved in the water. When the animals die or are eaten, the shells sink. Over many Calcium carbonate can be found in coral. centuries, thick layers of the shells may build up and be covered with mud, sand, or other materials. The shells are squeezed together by the heavy pressure of other materials and water above them. As they are squeezed together, the layer is converted to limestone. If the limestone is squeezed even more, it can change into marble or travertine.

Isotopes Six naturally occurring isotopes of calcium exist: calcium-40, calcium-42, calcium-43, calcium-44, calcium-46, and calcium-48. Isotopes are two or more forms of an element. Isotopes differ from each other according to their mass number. The number written to the right of the element's name is the mass number. The mass number represents the number of protons plus neutrons in the nucleus of an atom of the element. The number of protons determines the element, but the number of neutrons in the atom of any one element can vary. Each variation is an isotope. Radioactive isotopes of calcium have also been made. A radioactive isotope is one that breaks apart and gives off some form of radiation. Radioactive isotopes are produced when very small particles are fired at atoms. These particles stick in the atoms and make them radioactive. Two radioactive isotopes of calcium are used in research and medicine. Calcium-45 is used to study how calcium behaves in many natural processes. For example, it can be used to see how various types of soil behave with different kinds of fertilizers.

The calcium-45 is used as a tracer in such studies. A tracer is a radioactive isotope whose presence in a system can easily be detected. The isotope is injected into the system at some point. Inside the system, the isotope gives off radiation. That radiation can be followed by detectors placed around the system. Calcium-45 can also be used as a tracer in the study of glassy materials, detergents, and water purification systems. Both calcium-45 and calcium-47 can be used to study how calcium is used in the body. A doctor may think that a person's body is not using calcium properly in making bones or regulating nerve messages. The doctor can use calcium-45 or calcium-47 to find out more about this problem. The radioactive isotope is injected into the person's bloodstream. Then its path can be followed by the radiation it gives off. The doctor can then tell if the calcium is being used normally in the body. Extraction Pure calcium metal can be made by the same method used by Davy. An electric current is passed through molten calcium chloride: There is not much demand for pure calcium.

Most calcium is used in the form of limestone, gypsum, or other minerals that can be mined directly from the earth. Shellfish build their shells from calcium dissolved in the water. Uses Calcium metal has relatively few uses. It is sometimes used as a "getter." A getter is a substance that removes unwanted chemicals from a system. Calcium is used as a getter in the manufacture of evacuated glass bulbs. Calcium is added to the Calcium carbonate, as seen in the White Cliffs of Dover, England. bulb while it is being made. It then combines with gases left in the glass in the final stages of manufacture. Calcium is also used as a getter in the production of certain metals, such as copper and steel. The calcium removes unwanted elements that would otherwise contaminate the metal. Calcium is also used to make alloys. An alloy is made by melting and mixing two or more metals. The mixture has properties different from those of the individual metals. An alloy of calcium and cerium is used in flints found in lighters (the elements that create sparks). Compounds The starting point for the manufacture of most calcium compounds is limestone. Limestone occurs naturally in large amounts in many parts of the world.

It is usually mined from open-pit quarries. A quarry is a large hole in the ground from which useful minerals are taken. Limestone is first heated to obtain lime, or calcium oxide (CaO): Lime is one of the most important chemicals in the world. It usually ranks in the top five chemicals produced in the United States. In 1996, about 19 billion kilograms (42 billion pounds or 21 million tons) of lime was produced in the United States. Lime is used in the production of metals. It is used during the manufacture of steel to remove unwanted sand, or silicon dioxide (SiO 2 ), present in iron ore: The product formed in this reaction, calcium silicate (CaSiO 3 ), is called slag. Another important use of lime is in pollution control. Many factories release harmful gases into the atmosphere through smokestacks. Lining a smokestack with lime allows some of these gases to be captured. The lime is known as a scrubber.

Lime captures one harmful gas, sulfur dioxide (SO 2 ), which is a contributor to acid rain (a form of precipitation that is significantly more acidic than neutral water, often produced as the result of industrial processes): Calcium sulfite (CaSO 3 ) is a solid that can be removed from the inside of the smokestack. "In the limelight" At one time, lime was used as a source of light in theaters. When lime is heated to a high temperature, it gives off an intense white light. Pots of hot lime were often used to line the front of the stage. The light the pots gave off helped the audience see the performers. As a result, the performers were said to be "in the limelight." That phrase is still in use today, but lime is no longer used as a source of light in theaters. Lime is also used in water purification and waste treatment plants. When water combines with water, it forms slaked lime, or calcium hydroxide (Ca(OH) 2 ): Slaked lime traps impurities present in the water as it forms. It carries the impurities with it as it sinks to the bottom of the tank.

Lime is used to make more than 150 different industrial chemicals. Some examples of these chemicals with their uses are: Milk is a good source of calcium.

calcium alginate: thickening agent in food products such as ice cream and cheese products; synthetic fibers calcium arsenate (Ca 3 (AsO 4 ) 2 ): insecticide calcium carbide (CaC 2 ): used to make acetylene gas (for use in acetylene torches for welding); manufacture of plastics calcium chloride (CaCl 2 ): ice removal and dust control of dust on dirt roads; conditioner for concrete; additive for canned tomatoes; provides body for automobile and truck tires calcium cyclamate (Ca(C 6 H 11 NHSO 4 ) 2 ): sweetening agent (cyclamate), no longer permitted for use because of suspected cancer-causing properties calcium gluconate (Ca(C 6 H 11 O 7 ) 2 ): food additive; vitamin pills calcium hypochlorite (Ca(OCl) 2 ): swimming pool disinfectant; bleaching agent; deodorant; algicide and fungicide (kills algae and fungi) calcium permanganate (Ca(MnO 4 ) 2 ): liquid rocket propellant; textile production; water sterilizing agent; dental procedures calcium phosphate (Ca 3 (PO 4 ) 2 ): supplement for animal feed;

 fertilizer; commercial production of dough and yeast products; manufacture of glass; dental products calcium phosphide (Ca 3 P 2 ): fireworks; rodenticide (kills rats); torpedoes; flares calcium stearate (Ca(C 18 H 35 O 2 ) 2 ): manufacture of wax crayons, cements, certain kinds of plastics, and cosmetics; food additive; production of water resistant materials; production of paints calcium tungstate (CaWO 4 ): luminous paints; fluorescent lights; X-ray studies in medicine Health effects Calcium is essential to both plant and animal life.

In humans, it makes up about two percent of body weight. About 99 percent of the calcium in a person's body is found in bones and teeth. Milk is a good source of calcium. The body uses calcium in a compound known as hydroxyapatite (Ca 10 (PO 4 ) 6 (OH) 2 ) to make bones and teeth hard and resistant to wear. Calcium has many other important functions in the human body. For example, it helps control the way the heart beats. An excess (too much) or deficiency (not enough) of calcium can change the rhythm of the heart and cause serious problems. Calcium also controls the function of other muscles and nerves.

Also read article about Calcium from Wikipedia User Contributions: It's Elemental - The Element Potassium It's Elemental Melting Point: 336.53 K (63.38°C or 146.08°F) Boiling Point: 1032 K (759°C or 1398°F) Density: 0.89 grams per cubic centimeter Phase at Room Temperature: Solid Element Classification: Metal Period Number: 4 Group Number: 1 Group Name: Alkali Metal What's in a name? From the English word potash. Potassium's chemical symbol comes from the Latin word for alkali, kalium. Say what? Potassium is pronounced as poh-TASS-ee-em. History and Uses: Although potassium is the eighth most abundant element on earth and comprises about 2.1% of the earth's crust, it is a very reactive element and is never found free in nature. Metallic potassium was first isolated by Sir Humphry Davy in 1807 through the electrolysis of molten caustic potash (KOH). A few months after discovering potassium, Davy used the same method to isolate sodium .

Potassium can be obtained from the minerals sylvite (KCl), carnallite (KCl·MgCl2·6H2O), langbeinite (K2Mg2(SO4)3) and polyhalite (K2Ca2Mg(SO4)4·2H2O). These minerals are often found in ancient lake and sea beds. Caustic potash, another important source of potassium, is primarily mined in Germany, New Mexico, California and Utah. Pure potassium is a soft, waxy metal that can be easily cut with a knife. It reacts with oxygen to form potassium superoxide (KO2) and with water to form potassium hydroxide (KOH), hydrogen gas and heat. Enough heat is produced to ignite the hydrogen gas. To prevent it from reacting with the oxygen and water in the air, samples of metallic potassium are usually stored submerged in mineral oil. Potassium forms an alloy with sodium (NaK) that is used as a heat transfer medium in some types of nuclear reactors. Potassium forms many important compounds. Potassium chloride (KCl) is the most common potassium compound. It is used in fertilizers, as a salt substitute and to produce other chemicals.

Potassium hydroxide (KOH) is used to make soaps, detergents and drain cleaners. Potassium carbonate (KHCO3), also known as pearl ash, is used to make some types of glass and soaps and is obtained commercially as a byproduct of the production of ammonia. Potassium superoxide (KO2) can create oxygen from water vapor (H2O) and carbon dioxide (CO2) through the following reaction: 2KO2 + H2O + 2CO2 => 2KHCO3 + O2. It is used in respiratory equipment and is produced by burning potassium metal in dry air. Potassium nitrate (KNO3), also known as saltpeter or nitre, is used in fertilizers, match heads and pyrotechnics. Estimated Crustal Abundance: 2.09×104 milligrams per kilogram Estimated Oceanic Abundance: 3.99×102 milligrams per liter Number of Stable Isotopes: 2 ( View all isotope data ) Ionization Energy: 4.341 eV Who Discovered Calcium? Who Discovered Calcium?

Most people probably know that Calcium is part of the Periodic Table of the Elements with the other crucial substances. Perhaps, they also know other things about calcium such as: The atomic number is twenty, the symbol used in the table is Ca, and its atomic weight is at 40.078. But if someone would ask who discovered calcium, not many people would immediately have the answer for that. The Existence of Calcium Way Back Then According to history, calcium was already known during the early times – particularly, during the first century when Ancient Romans roamed and lived in the land. It was considered then as a particular component in the preparation of lime as an element called as calcium oxide. In those days, lime was called as calx. The Isolation Process and the Discovery of Calcium It was only during the year of 1808 did calcium get isolated and taken as a particular element in itself. And the person who discovered calcium was Englishman Sir Humphrey Davy. Davy was able to discover or, in effect, isolated this metallic element by combining mercuric oxide and lime and preparing the electrolysis of this combination. The temperature of the electrolysis of the mixture was done just a bit above the melting point.

This experiment was earlier influenced by the experiments done by Pontin and Berzelius. These two had a successful attempt in the preparation of calcium amalgam through the electrolyzing process of the lime substance in the substance of mercury. Learning about this success, Davy also made an attempt. And it was also a success, and led to the isolation or discovery of this soft and grayish element known as calcium. Through Davy’s life, this scientist was known to have a thorough knowledge in the process of electrolyzing substances and was actually the person who discovered or isolated other substances such as barium, magnesium, and strontium. The Properties of Calcium Calcium is classified as an alkaline earth metal that is a bit hard and is silvery white in color. Its melting point is around 839 +/-2 degrees C and boiling point is around 1484 degrees C. Its specification for gravity is at 1.55 or is equivalent to twenty degrees C, and its valence is two. Its Many Uses Calcium is also known to be in abundance, and is a crucial requirement for the nutrition of people. Even animals need this element since their skeletons need calcium phosphate. But not only people and animals thrive on this element.

Calcium is also very important for the growth and development of plants. It is also used in other metal preparations serving as a reducing agent, a decarbonizer, or as a scavenger. Fortunately, there was a person who discovered calcium, or, rather, had thought of separating this as an individual metallic element to further discover its various uses and components because now we know more about calcium and its value to people, animals, plants, and other elements. VN:F [1.9.22_1171] Potassium 16.4 x 106 S m-1 Freezing/Melting point: 63.4 oC, 336.5 K Davy (on the bellows) at a public demonstration of science at the Royal Institution in London. Image by James Gillray. Sodium and then potassium reacting with water. Discovery of Potassium Dr. Doug Stewart In 1806 English chemist Sir Humphry Davy discovered that chemical bonding was electrical in nature and that he could use electricity to split substances into their basic building blocks – the chemical elements. In 1807 he isolated potassium for the first time at the Royal Institution, London.

He electrolyzed dried potassium hydroxide (potash) which he had very slightly moistened by exposing it to the moist air in his laboratory. The electrolysis was powered by the combined output of three large batteries he had built. When he applied a voltage from his batteries to the potassium hydroxide, he found globules “having a high metallic lustre” collected at the negatively charged electrode. (1) Edmund Davy, who assisted in the experiment, described Sir Humphry Davy’s reaction to the production of potassium metal, his Eureka moment: “… when the minute globules of potassium burst through the crust of potash, and take fire as they entered the atmosphere, he could not contain his joy – he actually danced about the room in ecstatic delight; some little time was required for him to compose himself sufficiently to continue the experiment.” (2) Potassium was the first metal to be isolated by electrolysis. Davy was astonished at the new metal’s low density, observing that it floated on oil – something no other metal would do. He placed a piece of potassium in water and observed that the water, “decomposes it with great violence, an instantaneous explosion with brilliant flame.

He also (bravely) added potassium to hydrochloric acid and saw it burn with a bright red flame. (3) The name potassium is from the English word ‘potash’, originally meaning an alkali extracted with water in a pot of ash of burned wood or tree leaves. Potassium’s symbol K comes from ‘kalium’ the name of the element in Germany and Scandinavia. (4) Just a few days after isolating potassium, Davy isolated sodium for the first time using the same method. Interesting Facts about Potassium Potassium and its close periodic table neighbor sodium are solids at room temperature. Their alloys however are not. NaK alloys containing 40 to 90 percent of potassium by weight are liquids at room temperature. The commercially available 78% K, 22% Na alloy stays liquid at temperatures as low as -12.6 oC (9.3 oF). All living cells need potassium to maintain fluid balance, therefore we and all other forms of life on Earth need potassium minerals to survive. Potassium is available in all meats, plants and dairy products. Fruit and vegetables are the best sources of potassium. Several neurotoxins work by disrupting our cells’ biological use of potassium.

This can result in severe pain, or even death. These neurotoxins include agitoxin, charybdotoxin and margatoxin (scorpion stings), apamin (bee stings), and dendrotoxin (mamba snake bites). Most of the universe’s potassium atoms were made in the final moments of giant stars as they exploded in supernovae. Potassium is made in the oxygen burning shell of stars when they explode. This is not normal burning, of course; it is nuclear fusion. Potassium is made, along with several other elements including sulfur, and silicon, during explosive oxygen burning in supernovae. All plants need potassium to survive; over 90% of all human use of potassium compounds is in the manufacture of plant fertilizers. People whose diets are low in potassium can suffer from hypokalemia. Severe hypokalemia can be life threatening. Symptoms include irregular heartbeat, fatigue, muscle cramps and constipation. It is unusual for people to be deficient in potassium solely as a result of getting too little in their diets. Usually hypokalemia is caused by other issues such as diarrhea and/or vomiting, use of antibiotics, and kidney disease.

Most people are familiar with carbon dating, which uses the decay of the radioactive carbon-14 isotope to find the ages of once-living things such as animal and plant matter. The radioactive isotope potassium-40 gives us a way of dating rocks. Potassium-40 decays to argon-40 and calcium-40 with a half-life of 1.25 billion years. The ratio of potassium-40 to argon-40 trapped in rock is used to determine how long it is since the rock has solidified. Will potassium metal react with ice? The still image you see above might not be the whole story. advertisement Sodium - The Environmental Literacy Council The Environmental Literacy Council Of having discovered sodium. – Edmund Clerihew Bentley, British novelist Sodium is a soft, highly reactive metal that melts at a very low temperature, lower than that of boiling water (97.5 C). It is light enough to float on water and soft enough to be cut easily with a knife. It also reacts violently with water, producing enough heat to make it melt and to give off hydrogen gas, which then ignites.

It was first isolated by the English chemist Humphry Davy in 1807, by electrolysis of molten salt (NaCl) into its constituent elements sodium and chlorine. Sodium is represented by the symbol Na, its atomic number is eleven, and atomic weight 22.9898. It has only one naturally occurring isotope, sodium-23. It is an alkali metal, occurring in the far left hand column of the periodic table. It gives up one electron from its outer shell to form a positive ion. Because it is so reactive, it is never found in its pure, elemental state in nature; it must be isolated artificially. Pure sodium quickly looses its silvery luster when exposed to air and becomes dull gray, due to the formation of sodium oxide on the surface. This is analogous to the rusting of iron, but proceeds much more quickly. Sodium is the sixth most abundant element in the Earth’s crust. The salt water of the oceans contains 1.14 percent sodium; igneous rocks contain an average of 2.9 percent sodium.

Besides sodium chloride, other important sodium salts found in nature include sodium carbonate (soda), sodium borate (borax), sodium nitrate (Chile saltpeter), and sodium sulfate. A number of other sodium compounds are prepared from sodium chloride, such as sodium hydroxide (NaOH), also commonly known as caustic soda or lye. Lye is one of the ingredients in soap. Sodium bicarbonate, or baking soda, is also made from salt, and has a number of common uses. It makes bread rise, it can be used as an antacid (because it is a base, it neutralizes stomach acid), and it is used in fire extinguishers to generate carbon dioxide, which then displaces oxygen and suffocates the fire. Sodium plays an important role in human physiology. The concentration of Na+ ions inside cells is generally lower than it is in the extracellular fluids. This sodium ion gradient is partially counterbalanced by an opposite potassium ion (K+) gradient. The membranes of our cells contain an enzyme that acts as a “sodium pump,” maintaining these gradients of both Na+ and K+ ions across the membrane.

The electrical potential difference, or voltage, associated with these gradients is important for nerve transmission and muscle contraction. Sodium ion gradients are also responsible for various transport processes, including sugar transport in the intestine and amino acid transport in red blood cells. Sodium is lost along with moisture when we perspire (sweat is salty), so it is important to take in salt as well as water when exerting oneself in hot weather, to avoid dehydration and maintain the correct composition of body fluids. Pure sodium is produced by electrolysis of molten sodium chloride, which in turn is obtained from salt water. Calcium chloride is added to the sodium chloride to lower its melting point, from 800 C to 580 C. The lower temperature makes for less severe operating conditions, which then permits a simpler construction of the apparatus. Molten sodium is used as the heat-transfer medium in nuclear reactors. It is well suited to this purpose because it has low density (slightly less than that of water), low vapor pressure even at high temperatures, low viscosity, and very high heat capacity and heat conductivity.

Another use of pure sodium is in the reduction of animal and vegetable oils to long-chain fatty alcohols, which then serve as raw materials for detergents. Street lamps are commonly filled with sodium vapor, which accounts for their yellowish tint: sodium emits light of a very specific frequency when it is burned or excited by an electric discharge. The radioactive isotope sodium-24 is used as a radiotracer in biological research. Because this isotope has chemical behavior identical to that of common sodium, it forms the same compounds. Salt made with sodium-24 is taken up by the body and its progress through various physiological systems can be traced by sensors that detect the radiation given off. Because sodium-24 has a very short half-life (about 15 hours), radiation damage to the body is minimized. Recommended Resources The Royal Society of Chemistry: Chemical Science Network: Sodium The Chemical Science Network, provided by the Royal Society of Chemistry, the largest European organization in the chemical sciences, provides an interactive periodic table with information on each of the elements. The sodium page contains information about its discovery, its uses, and its biological role for humans and animals.

Chem4Kids: Sodium Andrew Rader, a University of California at Santa Barbara graduate with a background in science and computers, maintains the Chem4Kids site as a web-based tool for educators to teach chemistry basics. The page on sodium provides a description of the electrons in its shell and its atomic number, and how they determine its position on the periodic table. Rader also includes information on where one finds chlorine in every day life and the compounds it forms with other elements. The Jefferson Lab is managed and operated by Southeastern Universities Research Association for the U.S. Department of Energy. This page provides the basic physical and historical information about sodium, such as its atomic number, its uses, and its oxidation state. Related Content Potassium»historical information [WebElements Periodic Table] Potassium: historical information Potassium: historical information Potassium was discovered by Sir Humphrey Davy at 1807 in England. Origin of name : from the English word "potash" (pot ashes) and the Arabic word "qali" meaning alkali (the origin of the symbol K comes from the Latin word "kalium") Until the 18th century no distinction was made between potassium and sodium.

This was because early chemists did not recognise that "vegetable alkali" (K2CO3, potassium carbonate, coming from deposits in the earth) and "mineral alkali" (Na2CO3, sodium carbonate, derived from wood ashes) are distinct from each other. Eventually a distinction was made. Well before potassium was recognized as an element, potassium carbonate was mixed with animal fat to make soap. The carbonate was made by extracting wood ash with water before concentration by boiling - hence the name "potash" for potassium salts. Potassium was isolated in 1807 by Sir Humphry Davy, who obtained it through the electrolysis of very dry molten caustic potash (KOH, potassium hydroxide). Potassium collected at the cathode. Potassium was the first metal isolated by electrolysis. Davy isolated sodium by a similar procedure later in 1807. Sometime prior to the autumn of 1803, the Englishman John Dalton was able to explain the results of some of his studies by assuming that matter is composed of atoms and that all samples of any given compound consist of the same combination of these atoms.

Dalton also noted that in series of compounds, the ratios of the masses of the second element that combine with a given weight of the first element can be reduced to small whole numbers (the law of multiple proportions). This was further evidence for atoms. Dalton's theory of atoms was published by Thomas Thomson in the 3rd edition of his System of Chemistry in 1807 and in a paper about strontium oxalates published in the Philosophical Transactions. Dalton published these ideas himself in the following year in the New System of Chemical Philosophy. The symbol used by Dalton for potassium is shown below. [See History of Chemistry, Sir Edward Thorpe, volume 1, Watts & Co, London, 1914.] WebElements chemistry shop You can buy periodic table posters, mugs, T-shirts, periodic table fridge magnets, games, molecular models, and more at the WebElements periodic table shop potassium (K) | chemical element | Britannica.com chemical element hydrogen (H) Potassium ( K ), chemical element of Group 1 (Ia) of the periodic table , the alkali metal group, indispensable for both plant and animal life.

Potassium was the first metal to be isolated by electrolysis , by the English chemist Sir Humphry Davy , when he obtained the element (1807) by decomposing molten potassium hydroxide (KOH) with a voltaic battery . Four fragments of Potassium metal. Dennis S.K Encyclopædia Britannica, Inc. Properties, occurrence, and uses Potassium metal is soft and white with a silvery lustre, has a low melting point , and is a good conductor of heat and electricity. Potassium imparts a lavender colour to a flame, and its vapour is green. It is the seventh most abundant element in Earth’s crust, constituting 2.6 percent of its mass. Potash mine at Esterhazy, Sask. George Hunter/Shostal Associates The potassium content of the Dead Sea is estimated at approximately 1.7 percent potassium chloride, and many other salty bodies of water are rich in potassium. The waste liquors from certain saltworks may contain up to 40 grams per litre of potassium chloride and are used as a source of potassium.

Most potassium is present in igneous rocks, shale, and sediment in minerals such as muscovite and orthoclase feldspar that are insoluble in water; this makes potassium difficult to obtain. As a result, most commercial potassium compounds (often loosely called potash) are obtained via electrolysis from soluble potassium compounds, such as carnallite (KMgCl3∙6H2O), sylvite (potassium chloride, KCl), polyhalite (K2 Ca 2Mg[SO4]4∙2H2O), and langbeinite (K2 Mg 2[SO4]3), which are found in ancient lake beds and seabeds. alkali metal: Potassium is produced by sodium reduction of molten potassium chloride, KCl, at 870 °C (1,600 °F). Molten KCl is continuously fed into a packed distillation column while sodium vapour is passed up through the column. By condensation of the more volatile potassium at the top of the distillation tower, the reaction Na + KCl → K + NaCl is forced to the right.

Efforts to devise a scheme for commercial electrolytic production of potassium have been unsuccessful because there are few salt additives that can reduce the melting point of potassium chloride to temperatures where electrolysis is efficient. Similar Topics rhodium (Rh) There is little commercial demand for potassium metal itself, and most of it is converted by direct combustion in dry air to potassium superoxide, KO2, which is used in respiratory equipment because it liberates oxygen and removes carbon dioxide and water vapour. (The superoxide of potassium is a yellow solid consisting of K+and O2− ions. It also can be formed by oxidation of potassium amalgam with dry air or oxygen.) The metal is also used as an alloy with sodium as a liquid metallic heat-transfer medium. Potassium reacts very vigorously with water, liberating hydrogen (which ignites) and forming a solution of potassium hydroxide, KOH. Sodium-potassium alloy (NaK) is used to a limited extent as a heat-transfer coolant in some fast-breeder nuclear reactors and experimentally in gas-turbine power plants. The alloy is also used as a catalyst or reducing agent in organic synthesis. In addition to the alloys of potassium with lithium and sodium, alloys with other alkali metals are known.

Complete miscibility exists in the potassium-rubidium and potassium-cesium binary systems. The latter system forms an alloy melting at approximately −38 °C (−36 °F). Modification of the system by the addition of sodium results in a ternary eutectic melting at approximately −78 °C (−108 °F). The composition of this alloy is 3 percent sodium, 24 percent potassium, and 73 percent cesium . Potassium is essentially immiscible with all the alkaline-earth metals, as well as with zinc , aluminum , and cadmium . Potassium (as K+) is required by all plants and animals. Plants need it for photosynthesis, regulation of osmosis and growth, and enzyme activation. Every animal has a closely maintained potassium level and a relatively fixed potassium-sodium ratio. Potassium is the primary inorganic cation within the living cell, and sodium is the most abundant cation in extracellular fluids. In higher animals, selective complexants for Na+ and K+ act at cell membranes to provide “active transport.” This active transport transmits electrochemical impulses in nerve and muscle fibres and in balancing the activity of nutrient intake and waste removal from cells.

Too little or too much potassium in the body is fatal, but potassium in the soil ensures the presence of this indispensable element in food. Britannica Stories Big Radio Burst from Tiny Galaxy The potassium content of plants varies considerably, but it is ordinarily in the range of 0.5–2 percent of the dry weight. In humans the ratio of potassium between the cell and plasma is approximately 27:1. The potassium content of muscle tissue is approximately 0.3 percent, whereas that of blood serum is about 0.01–0.02 percent. The dietary requirement for normal growth is approximately 3.3 grams (0.12 ounce) of potassium per day, but the ingestion of more than 20 grams (0.7 ounce) of potassium results in distinct physiological effects. Excess potassium is excreted in the urine, and a significant quantity may be lost during sweating. Periodic Table of the Elements Natural potassium consists of three isotopes : potassium-39 (93.26 percent), potassium-41 (6.73 percent), and radioactive potassium-40 (about 0.01 percent); several artificial isotopes have also been prepared.

Potassium-39 is normally about 13.5 times more plentiful than potassium-41. The natural radioactivity of potassium is due to beta radiation from the potassium-40 isotope (109 years half-life). The disintegration of potassium-40 is used in geological age calculations (see potassium-argon dating ). Potassium easily loses the single 4s electron , so it normally has an oxidation state of +1 in its compounds, although compounds that contain the anion , K-, can also be made. Principal compounds and reactions with other elements Of commercially produced potassium compounds, almost 95 percent of them are used in agriculture as fertilizer . (Potassium compounds are also important to a lesser extent in the manufacture of explosives.) The world supply of potash for fertilizer is about 25 million tons (calculated as K2O, although potassium in fertilizer is most commonly present as KCl). Large deposits of sylvite in Saskatchewan , Can., provide more than 25 percent of the world’s needs. The other chief sources of potash are Germany , Russia , Belarus , India, Chile, and Israel . Seawater , brines, and ashes of vegetation are also used as sources of potash.

Connect with Britannica Facebook Twitter YouTube Instagram Pinterest Potassium chloride , KCl, is a naturally occurring potassium salt that, aside from its use as fertilizer, is also a raw material for the production of other important potassium compounds. Electrolysis of potassium chloride yields potassium hydroxide (also called caustic potash ), which readily absorbs moisture and is employed in making liquid soaps and detergents and in preparing many potassium salts . Reaction of iodine and potassium hydroxide produces potassium iodide , KI, which is added to table salt and animal feed to protect against iodine deficiency. Other potassium compounds of economic value include potassium nitrate, also known as saltpetre , or nitre, KNO3, which has wide use as a fertilizer and in fireworks and explosives and has been used as a food preservative; potassium chromate , K2CrO4, which is employed in tanning leather and dyeing textiles; and potassium sulfate, K2SO4, which is used in the production of fertilizers and potassium alums . The chemical properties of potassium are similar to those of sodium, although the former is considerably more reactive. Potassium differs from sodium in a number of respects.

Whereas sodium is essentially unreactive with graphite, potassium reacts to form a series of interlamellar compounds, the richest having the formula KC8. Compounds are formed with carbon–potassium atomic ratios of 8, 16, 24, 36, 48, and 60 to 1. The graphite lattice is expanded during penetration of the potassium between the layers. Potassium reacts with carbon monoxide at temperatures as low as 60 °C (140 °F) to form an explosive carbonyl (K6C6O6), a derivative of hexahydroxybenzene. Liquid potassium and NaK both are more reactive than liquid sodium with air and oxygen. Potassium reacts violently with water to produce half a mole of hydrogen per mole of potassium and water and generates approximately 47 kilocalories per mole of heat. Potassium can be stored in nitrogen gas with no reaction. It reacts with hydrogen at approximately 350 °C (660 °F) to form the hydride. Potassium is highly reactive with halogens and detonates when it contacts liquid bromine .

Violent explosions also have been observed when mixtures of potassium and halogen acids are subject to shock. Explosions also have occurred when potassium is mixed with a number of metal halide salts or with organic-halogen compounds. Britannica Lists & Quizzes

Question: Who discovered the chemical elements Potassium, Sodium and Calcium?
Answer: HUMPHREY DAVY